General Chemistry Study Notes: Reactions in Aqueous Solution, Thermochemistry, and Electronic Structure

Reactions in Aqueous Solution

Solubility and Precipitation Reactions

Solubility rules help predict whether a compound will dissolve in water or form a precipitate. Precipitation reactions occur when two aqueous solutions are mixed and an insoluble product (precipitate) forms.

• Ammonium Ion (NH4+): Generally forms soluble salts, but exceptions exist.

• Common Insoluble Salts: Carbonates (except with alkali metals and ammonium), sulfates (some exceptions), and chlorides (except Ag+, Pb2+, Hg22+).

• Example: Mixing solutions of KCl and AgNO3 forms AgCl precipitate.

Acids and Bases: Diprotic Acids

Acids are classified by the number of protons (H+) they can donate. A diprotic acid can donate two protons per molecule.

• Diprotic Acid: An acid that can release two H+ ions (e.g., sulfuric acid, H2SO4).

• Monoprotic Acid: Can donate only one proton (e.g., nitric acid, HNO3).

• Example: H2SO4 → 2H+ + SO42−

Redox Reactions: Oxidation and Reduction

Redox reactions involve the transfer of electrons between substances. Oxidation is the loss of electrons, while reduction is the gain of electrons.

• Oxidation: Loss of electrons (may also involve gain of oxygen).

• Reduction: Gain of electrons (may also involve loss of oxygen).

• Mnemonic: "OIL RIG" – Oxidation Is Loss, Reduction Is Gain (of electrons).

• Example: Zn → Zn2+ + 2e− (oxidation); Cu2+ + 2e− → Cu (reduction)

Acid-Base Neutralization

When an acid reacts with a metal hydroxide (base), a neutralization reaction occurs, producing water and a salt.

• General Equation: $\text{Acid} + \text{Base} \rightarrow \text{Salt} + \text{Water}$

• Example: HCl + NaOH → NaCl + H2O

Strong Electrolytes

Electrolytes are substances that dissociate into ions in solution. Strong electrolytes completely ionize in water.

• Strong Electrolyte: Completely ionizes in solution (e.g., NaCl, HCl).

• Weak Electrolyte: Partially ionizes (e.g., acetic acid).

• Nonelectrolyte: Does not ionize (e.g., sugar).

Spectator Ions

Spectator ions are ions that do not participate in the chemical reaction and remain unchanged in solution.

• Example: In the reaction KCl (aq) + AgNO3 (aq) → AgCl (s) + KNO3 (aq), K+ and NO3− are spectator ions.

Dilution Calculations

When a solution is diluted, the amount of solute remains constant, but the concentration decreases as the volume increases.

• Formula: $M_1V_1 = M_2V_2$

• Example: 400 mL of solution diluted to 4.00 L, final concentration 0.0400 M. Initial concentration: $M_1 = \frac{M_2V_2}{V_1}$

Calculating Mass from Molarity

To find the mass of solute in a solution, use the molarity, volume, and molar mass.

• Formula: $\text{mass} = M \times V \times \text{molar mass}$

• Example: 500.0 mL of 0.250 M NaOH: $\text{mass} = 0.250 \times 0.500 \times 40.0 = 5.00$ g

Thermochemistry

Endothermic and Exothermic Reactions

Thermochemistry studies the heat changes in chemical reactions. Reactions can absorb or release heat.

• Endothermic Reaction: Absorbs heat from surroundings; $\Delta H$ is positive.

• Exothermic Reaction: Releases heat to surroundings; $\Delta H$ is negative.

• Example: Melting ice (endothermic), combustion (exothermic).

Identifying Exothermic Processes

Exothermic processes release energy, often as heat, to the surroundings.

• Examples:

  • Condensation of water vapor

  • Boiling soup (if referring to heat release from chemical change)

  • Combustion reactions

• Endothermic Examples: Ice melting, water evaporating

Electronic Structure of Atoms (Additional info)

Oxidation and Reduction: Electron Transfer

Electron transfer is central to redox reactions, which are important in energy production and corrosion.

• Oxidation Number: Indicates the degree of oxidation of an atom in a compound.

• Example: In NaCl, Na has +1, Cl has -1 oxidation state.